Introduction to Halides group:

We are familiar that group 15 elements form compounds mostly in two oxidation states, +3 and +5. Both kinds of their halides, that is, trihalides and also pentahalides are identified.

Trihalides:

All the elements of Group 15 form trihalides, MX₃. Of all nitrogen halides, just NF₃ is stable because of the strong N-F bonds. All the other trihalides of nitrogen, NCl₃, NBr₃ and NI₃ are unstable. Of such, NCl₃ is explosive and NBr₃ and NI₃ exist just as their ammoniates, that is, as ammonia adducts and detonate on eliminating ammonia.

Apart from NX₃, all the trihalides hydrolyze quickly to either hydrated oxides, example -As2O3·nH2O or acids, H₃PO₃ or oxochloride, SbOCl and BiOCl. NCl₃ on hydrolysis provides NH₃ and HOCl rather than HNO₂ (nitrous acid) and HCl:

NCl₃ + 3H₂O → NH₃ + 3HOCl 

PCl₃ + 3H₂O → H₃PO₃ + 3HCl

SbCl₃ + H₂O → SbOCl + 2HCl

BiCl₃ + H₂O → BiOCl + 12HCl

The method or procedure of the hydrolysis of phosphorus trichloride is identical to that of silicon tetrachloride. It comprises the formation of an intermediate four coordinate species H₂O→PCl₃. The formation of this kind of intermediate is not possible for nitrogen due to the non-availability of d-orbitals. NCl₃, thus, hydrolyses via a different method, which comprises the attack of the Lewis base, water, on the partially positive chlorine to eliminate it in the form of a protonated hypochlorous acid molecule. Proton exchange by the Cl₂N^- anion then produces one molecule of HOCl and NHCl₂ each.

Substitution of two more chlorine atoms by hydrogen produces the NH3 molecule:

Fig: Hydrolysis of Nitrogen trichloride

The trihalides, apart from NX₃, act as Lewis acids, that is, electron pair acceptors, forming halo anions, SbCl₄, BiCl₄, and so on. The trihalides can be oxidized to pentahalides. The easiness of oxidation reduces from PX₃ to BiX₃. Though, NX₃ is not oxidized at all. The trihalides function as halogenating agents. PCl₃ reacts by NH₃, halogens, HI, S, O₂ and AgNCO to provide P(NH₂)₃, PX₂Cl₃, PI₃, PSCl₃, POCl₃ and  P(NCO)₃, correspondingly. The given chemical equations symbolize these reactions:

PCl₃ + 3NH₃ → P(NH₂)₃ + 3HCl

PCl₃ + Br₂ → PBr₂Cl₃

PCl₃ + HI → PI₃ + 3HCl 

PCl₃ + S → PSCl₃ 

2PCl₃ + O₂ → 2POCl₃

PCl₃ + 3AgNCO → P(NCO)₃ + 3AgCl

Trihalides of the lighter elements of the group are mostly covalent in nature, thus, they can exist as discrete molecules that encompass a pyramidal structure, such as NH₃. The trihalides of heavier elements, example - BiF₃, are ionic in nature.

Pentahalides:

Nitrogen doesn't form any pentahalides. The stability of other pentahalides of the group reduces in the order P > Sb > As > Bi and in the order F > Cl > Br for halogens. PF₅, PCl₅, PBr₅, AsF₅, SbF₅, SbCl₅, and BiF5 are stable. Though, PBr₅ and SbCl₅ readily lose halogen forming trihalides. The pentahalides get more oxidising as we go down in the group, apart from arsenic pentahalides, AsX₅, which are unpredictably more oxidizing than those of the element following it in the group. Each and every pentahalides encompass trigonal bipyramidal structure in the gas phase. Though, their structures vary in the solid state, example - PCl₅ divides into [PCl₄] [PCl₆] ¯ which encompass tetrahedral and octahedral structures, correspondingly; PBr₅ divides to give [PBr₄]⁺Br^-. Pentahalides are obtained via the action of surplus of halogen on trihalides:

PX₃ + X₂ ↔ PX₅

Pentahalides are good halogenating agents, example - PCl₅ is employed for conversion of alcohols to alkyl halides and acids to acylchlorides as illustrated below:

ROH + PCl₅ → RCl + POCl₅ + HCl

RCOOH + PCl₅ → RCOCl + POCl₃ + HCl

They hydrolyze to give oxoacids, H₃PO₄, H₃AsO₄, or hydrated oxides such as Sb₂O, nH₂O or oxohalides such as BiOCl; partial hydrolysis of PCl₅ gives POCl₃. Nearly all the pentahalides can accept the electron pair and act as Lewis acids, as in PF₆, PF₅ and so on.

Introduction of Oxides group:

As stated, nitrogen forms a number of oxides, N₂O, NO, N₂O₃, NO₂ or N₂O₄ and N₂O₅, and as well very not stable NO₃ and N₂O₆. All such oxides of nitrogen show pπ-pπ multiple bonding between nitrogen and oxygen. This doesn't take place with the heavier elements in the group and as a result nitrogen forms a number of oxides which encompass no P, As, Sb or Bi analogues.

Oxides of Nitrogen:

Significant features of oxides of nitrogen are illustrated in the table shown below. Let us first illustrate the preparation, properties and structures of these oxides.

Table: Oxides of Nitrogen

=> Preparation:

N₂O is obtained usually by heating NH₄NO₃:

NH₄NO₃ → N₂O + 2H₂O

NO is prepared by the reduction of 8M HNO₃ by reducing agents such as Cu or by reduction of nitrous acid or nitrites by Fe²⁺ or I^- ions:

3Cu + 8HNO₃ → 3Cu(NO₃)₂ + 2NO + 4H₂O

2NaNO₂ + 2FeSO₄ + 3H₂SO₄ → 2NaHSO₄ + Fe₂(SO₄)₃ + 2NO + 2H₂O

2NaNO₂ + 2NaI + 4H₂SO₄ → 4NaHSO₄ + 2NO + I₂ + 2H₂O

NO is formed as an intermediate in the preparation of nitric acid via oxidation of NH₃. N₂O₃ is obtained as an intense blue liquid or a pale blue solid on cooling the equimolar mixture of NO and NO₂:

NO + NO2 → N₂O₃

On warming, its colour fades because of its dissociation into these two oxides.

NO₂ can be manufactured by reduction of cone. HNO₃ by Cu or through heating heavy metal nitrates:

Cu + 4HNO₃ → Cu(NO₃)₂ + 2NO₂ + 2H₂O

2Pb(NO₃)₂ → 2PbO + 4NO₂ + O₂

N₂O₅ is the anhydride of HNO₃. This is best made by dehydrating HNO₃ by P₄O₁₀ at low temperatures:

4HNO₃ + P₄O₁₀ → at 250K → 2N₂O₅ + 4HPO₃

=> Properties and Structure:

The oxides of nitrogen are all oxidising agents, N₂O even supporting the combustion of S and P. NO that is thermally more stable, supports the combustion of Mg and P however not of S. Sulphur flame is not hot adequate to decompose it. N₂O and NO are neutral, whereas the other oxides are acidic.

N₂O is isoelectronic by CO2 and as well consists of a linear structure. Though, dissimilar to CO₂, N₂O consists of a small dipole moment. The resonance structures of the molecule are given in the figure shown below.

Fig: Resonance Structures of Nitrous Oxide

NO consists of a total of 15 electrons. This is impossible for all of them to be paired and therefore this is an odd electron molecule. In the gaseous state, it is paramagnetic. Though, the liquid and solid states are diamagnetic as loose dimers (figure shown below) are formed cancelling out the magnetic effect of the unpaired electrons.

Fig: Resonance Structures of Nitric Oxide

The bonding in NO is best illustrated by the molecular orbital theory. Molecular orbital electronic configuration of NO molecule can be symbolized as σ1s², σ*1s², σ2s², σ*2s², σ(2p_z)², π(2p_x)², π(2p_y)², π*(2p_x)¹. This provides a bond order of 2.5. If the unpaired electron occupying the anti-bonding π*2p_y orbital is removed, nitrosonium ion, NO⁺, is formed and the bond order becomes 3. This is reflected in the shortening of the bond length from 115 pm in NO to 106 pm in NO⁺. Nitrosonium ion is stable and forms salts such as NO⁺Cl¯. This is isoelectronic with CO and forms complexes with the transition elements. The brown ring made up in the test for nitrates is due to the formation of a complex of iron, [Fe(H₂O)5NO]²⁺.

NO₂ having 23 electrons is again an odd electron molecule. In the gaseous state it is paramagnetic. On cooling, the gas condenses to a brown liquid and ultimately to a colourless solid both of which are diamagnetic because of dimerisation. NO₂ molecule is angular by ONO angle of 134°. The O-N bond length is 120 pm, intermediate between the single and double bond. The odd electron is on nitrogen. The dimer has been illustrated to encompass a planar structure. The N-N bond length is extremely large, 175 pm, making this an extremely weak bond.

Fig: Resonance Structure of (a) NO₂ and (b) N₂O₄

Oxides of Phosphorus, Arsenic, Antimony and Bismuth: 

As P, As and Sb form oxides in both +3 and +5 oxidation states while only one oxide of bismuth, Bi₂O₃ is known. Keep in mind that the stability of higher oxidation states decreases on going down the group. Apart from Bi₂O₃ others exist in dimeric form. The fundamental character of oxides increases on descending the group. The oxides of P and As are acidic, those of Sb amphoteric and of Bi completely basic. As well the higher oxidation states are more acidic.

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